Understanding Periodicity: Covalent Radius, Ionisation Energy and Electronegativity in Chemistry

Covalent Radius

The covalent radius is half the distance between the nuclei of two covalently bonded atoms. This measurement helps us understand atomic size trends.

Across a period, covalent radius decreases. This happens because as you move right across the periodic table, more protons are added to the nucleus while electrons go into the same shell. The stronger nuclear pull draws electrons closer to the nucleus, shrinking the atom.

Down a group, covalent radius increases. Each element down a group has an extra electron shell, making the atom physically larger. Inner electron shells also create a shielding effect that reduces the nuclear attraction on outer electrons.

Quick Tip: Remember the pattern - atoms get smaller as you move right across the periodic table, but larger as you move down a group.

First Ionisation Energy

First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms. For example, with magnesium: Mg$g$ → Mg⁺$g$ + e⁻. This is always an endothermic process measured in kJ/mol.

Across a period, ionisation energy increases. The growing nuclear charge creates stronger attraction to electrons, making them harder to remove.

Ionisation Energy Trends

Down a group, ionisation energy decreases. Outer electrons are further from the nucleus and shielded by inner electron shells. This reduced attraction means less energy is needed to remove an electron.

The second ionisation energy involves removing another electron from an already positive ion. For example: Mg⁺$g$ → Mg²⁺$g$ + e⁻. This always requires more energy than the first ionisation energy because you're removing an electron from a positively charged ion where the remaining electrons are held more tightly.

To find the total energy needed to remove two electrons, you add the first and second ionisation energies together.

Remember: Second ionisation energies are always higher than first ionisation energies because you're removing an electron from a positively charged ion where electrons are less shielded and more strongly attracted to the nucleus.

Electronegativity

Electronegativity measures how strongly an atom attracts electrons in a chemical bond. This property helps predict bond polarity and molecular behaviour.

Across a period, electronegativity increases. As nuclear charge grows and atomic size shrinks, atoms develop a stronger pull on bonding electrons. This explains why non-metals (right side of the periodic table) typically form anions by gaining electrons.

Down a group, electronegativity decreases. Despite increasing nuclear charge, the greater number of electron shells creates a significant shielding effect. With bonding electrons kept further from the nucleus, the atom's electron-attracting ability weakens.

Exam Tip: Electronegativity trends follow the same pattern as ionisation energy - increasing across periods and decreasing down groups. This makes sense as both relate to an atom's ability to control electrons.